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Notes Chapter 5 “The Periodic Law” Section 1 Dmitri Mendeleev – father of the modern periodic table. Listed elements by masses and properties. Left a column for the noble gases. Publishes in 1869. Henry Mosely working with Rutherford in 1911 – helped come up with the “periodic law” – physical and chemical properties are functions of the atomic number. Today periodic table is arranged according to atomic number and properties. 1894 Strutt and Ramsay discovered Argon the first of the noble gases. The Lanthanide – All end in 4f, numbers 58-71, period 6. Actinides – all end in 5f, number 90-103. Belongs in period 7. **to save space the lanthanides and actinides are set off below the main table. 8, 8, 18,18, 32 –periodicity Section 2 See table 1 and figure 5. S block – groups 1 and 2 – group 1(alkali metals- end in s1, not found in nature as free elements. Always found in compounds. React strongly with water. Soft silvery appearance can be cut with a knife, will have a +1 charge). Group 2( alkaline earth metals, end in s2, harder, denser, and stronger than group 1, not as reactive as group 1, will have a +2 charge). Hydrogen is not an alkali metal, it is located above the group because of it electron configuration. Helium possesses special chemical stability, and is unreactive so it is above group 18. D-block – transition elements all end in d. Metals with typical metallic properties. They are good conductors, electricity, and have luster. Less reactive than group 1 and 2. They do not easily form compounds. Will have a + charge. P block – all end with P, mixed group. Metals form + ions and nonmetals form – ions. Group 17(halogens) most reactive group in the p block, have -1 charge. Main- group elements- s and p block. Section 3 Atomic Radii- half the distance between the nuclei of atoms that are bonded together. Atomic radii decreases from left to right and across a period and increase down a group. Ionization energy – energy required to remove an electron. Ionization increase as electrons are removes ex. 1st electron removed has the lowest ionization energy. See figure 15 Electron Affinity – energy required to acquire and electron. Most atoms release energy when they acquire an electron. Some atoms have to be forced to gain electrons, this will require energy. See figure 17 Ionic Radii – *cation – positive, usually metals, smaller radius, because electrons have been removed. *anion – negative, usually nonmetals, larger radius, because electrons have been gained. Valence electrons – outer most s and p; gained, lost, or shared to form bonds. The P must follow the S. Electronegativity – atoms ability to attract electrons from another atom in a compound. See figure 20. Tends to increase across a period.